Organic Acids and Bases

Purpose

To identify the acid, base, conjugate acid, and conjugate base in representative organic Brønsted acid-base reactions; to use pKₐ values to predict the direction of equilibrium in each reaction; and to rank a set of organic O–H and C–H acids in order of increasing acidity based on structural features (electronegativity, hybridization, resonance, and inductive effects).

Theory

In Brønsted-Lowry theory, an acid donates a proton (H⁺) and a base accepts one. Every acid-base reaction is the transfer of a single proton between the two species, producing a conjugate acid-base pair on each side of the equilibrium: the acid (HA) becomes its conjugate base (A⁻), and the base (B) becomes its conjugate acid (BH⁺). The acid dissociation constant Kₐ (or its logarithmic form pKₐ = −log Kₐ) quantifies acid strength, with lower pKₐ corresponding to a stronger acid. An acid-base equilibrium favours the side with the weaker acid (higher pKₐ); quantitatively K_eq ≈ 10^(pKₐ_{product-side acid} − pKₐ_{reactant-side acid}).

The acidity of an H–X bond depends on the stability of the conjugate base A⁻. Four effects control this stability: (i) electronegativity of the atom bearing the charge (O⁻ more stable than N⁻ more stable than C⁻, within the same row); (ii) hybridization (sp lone pair is closer to the nucleus than sp³, so sp C–H is most acidic of the C–H family); (iii) resonance stabilisation of the conjugate base (carboxylate spreads charge across two oxygens, phenolate spreads charge onto the ring, enolate spreads charge onto oxygen); and (iv) inductive effects of substituents (EWGs like F, Cl, NO₂ stabilise the conjugate base, lowering pKₐ; EDGs like alkyl groups destabilise it, raising pKₐ).

Calculations / Worked Examples

Sample reaction — CH₃COOH + NH₃ ⇌ CH₃COO⁻ + NH₄⁺:

Acid (reactants): CH₃COOH, pKₐ = 4.76
Base (reactants): NH₃
Conjugate base (products): CH₃COO⁻
Conjugate acid (products): NH₄⁺, pKₐ = 9.25
ΔpKₐ = 9.25 − 4.76 = +4.49
K_eq ≈ 10^(+4.49) ≈ 3.1 × 10⁴
Equilibrium favours PRODUCTS (right) — acetic acid is the stronger acid, so it gives up its proton to NH₃ readily.

Sample ranking — O–H acids (increasing acidity):

tert-butanol (17) < ethanol (16) < water (15.7) < phenol (10) < p-nitrophenol (7.15) < CF₃COOH (0.23)
Reasoning: tert-butanol's three alkyl groups donate electron density inductively, destabilising its conjugate base (t-BuO⁻) and raising pKₐ above ethanol. Water is slightly more acidic than primary alcohols because the hydroxide conjugate base has less alkyl destabilisation. Phenol (pKₐ 10) is much more acidic than any aliphatic alcohol because the phenolate anion is resonance-stabilised by delocalisation onto the aromatic ring (3 additional resonance structures place negative charge on ortho/para carbons). Adding a para-nitro group to phenol further lowers pKₐ to 7.15: the nitro group is a strong EWG that stabilises the phenolate through both inductive pull and resonance with the ring. Trifluoroacetic acid's three fluorines pull electron density inductively from the carboxylate, stabilising it dramatically — CF₃COO⁻ is so stable that CF₃COOH has pKₐ ≈ 0, making it ~3 × 10⁴ times more acidic than acetic acid.

Results Table

Section I — component identification and equilibrium direction

Section II — acidity rankings

Discussion

Section I confirmed that the direction of every Brønsted equilibrium is predicted by the pKₐ difference between the acid on each side. Reactions where the reactant acid has a lower pKₐ than the conjugate acid formed on the product side (positive ΔpKₐ) proceed forward — CH₃COOH is a stronger acid than NH₄⁺, so it transfers its proton to NH₃ almost completely (K_eq ≈ 3 × 10⁴). Conversely, attempting to deprotonate a ketone α-H (pKₐ 19.3) with hydroxide (H₂O pKₐ 15.7) is unfavourable because the enolate is a stronger base than hydroxide — less than 1 in 10⁴ molecules is deprotonated at equilibrium, which is why kinetically-controlled enolate chemistry typically uses stronger bases like LDA (pKₐ of HN(i-Pr)₂ ≈ 36) rather than NaOH.

The α-H acidity of carbonyls is a clear illustration of resonance stabilisation: although acetone's α-C–H is formally an sp³ C–H bond that would normally be pKₐ ~50, the resulting enolate delocalises the negative charge onto the electronegative carbonyl oxygen, dropping the pKₐ to 19.3 — a stabilisation of over 30 pKₐ units (10³⁰ in equilibrium constant). Adding a second carbonyl (as in 2,4-pentanedione or diethyl malonate) delocalises the charge onto two oxygens, dropping pKₐ further to 9.0, comparable to phenol and ammonium.

Section II's rankings exposed the four factors driving acidity in a quantitative way. Round 1 (O–H acids) spans 17 pKₐ units (10¹⁷-fold difference in Kₐ), driven almost entirely by conjugate-base stabilisation. Aliphatic alkoxides sit at the top (least acidic) because alkyl groups destabilise the negative oxygen through weak electron donation. Phenoxide benefits from resonance into the ring, dropping pKₐ to 10; a p-nitro group reinforces this via both resonance and induction, reaching 7.15. Carboxylate's two-oxygen resonance drops acetic acid to 4.76, and triple-fluorine inductive stabilisation on CF₃COO⁻ reaches pKₐ ≈ 0. Round 2 (C–H acids) emphasised hybridization (sp 25 < sp² 43 < sp³ 50) and the dramatic effect of α-conjugation (acetone 19.3; acac 9.0).

Conclusion

All six reactions in Section I were labelled correctly and the directions predicted quantitatively via ΔpKₐ; four of six favoured products, two favoured reactants, matching the rule that equilibrium favours the weaker acid. The O–H and C–H acidity rankings in Section II reproduced the expected trends based on electronegativity, hybridization, resonance, and inductive effects, with the full range of organic acidity spanning nearly 50 pKₐ units. All four unknowns were identified from their pKₐ and structural clues. The experiment confirmed that conjugate-base stability, not intrinsic proton "affinity", is the organising principle of organic acid-base chemistry.